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JEE Chemistry
Line Spectrum of Hydrogen

Line Spectrum of Hydrogen

An emission spectrum shows the wavelengths emitted by a substance after absorbing energy, with excited atoms releasing energy as they return to a lower state. This results in a line spectrum in gases, where specific wavelengths appear as bright lines separated by dark gaps unique to each element. Conversely, an absorption spectrum occurs when a substance absorbs specific wavelengths from a continuous spectrum, creating dark lines at the absorbed wavelengths.

1.0Spectral Emission

Spectral Emission occurs when an electron moves from a higher energy level (n₂) to a lower energy level (n₁). This transition releases energy as a photon, with the photon's energy corresponding to the difference between the two energy states. Since the energy levels of an atom are quantised and fixed, the emitted photon will always have the same energy for a given transition. This leads to a spectrum with specific wavelengths that reflect the differences between the energy levels in the atom.

2.0Spectral Series of Hydrogen

As the simplest atomic system, the hydrogen atom exhibits the most straightforward spectral series. When a beam of light or radiation passes through a slit and is observed using a spectroscope, it creates a series of images that appear as parallel lines. These lines are arranged side-by-side with uniform spacing, representing individual components of the radiation. As one moves from higher to lower wavelengths, the spectral lines are more widely spaced on the high-wavelength side and progressively closer together on the low-wavelength side. The shortest wavelength, the series limit, features the most closely packed spectral lines.

Spectral Lines

In 1885, Johann Jakob Balmer, a Swedish teacher, first identified a series in the visible region of the hydrogen spectrum known as the Balmer series. This series corresponds to electron transitions from higher energy levels to the second energy level (shell).

Other notable hydrogen spectral series include:

  • Lyman Series: Electron transitions to the first energy level from any higher level.
  • Balmer Series: Electron transitions to the second energy level from any higher level.
  • Paschen Series: Electron transitions to the third energy level from any higher level.
  • Bracket Series: Electron transitions to the fourth energy level from any higher level.
  • Pfund Series: Electron transitions to the fifth energy level from any higher level.

hydrogen spectral series

Each series corresponds to specific regions of the electromagnetic spectrum, ranging from ultraviolet (Lyman) to infrared (Paschen, Bracket, and Pfund).

Spectral Series

Spectral Region

n₁

n₂

Wavelength (Å)

Lyman Series

Ultraviolet

1

2, 3, 4, 5, 6, 7, …..

920 - 1200

Balmer Series

Visible light

2

3, 4, 5, 6, 7, ...

4000 - 6500

Paschen Series

Infrared

3

4, 5, 6, 7, ...

9500 - 18750

Brackett Series

Infrared

4

5, 6, 7, ...

19450 - 40500

Pfund Series

Infrared

5

6, 7, ...

37800 - 75000

Humphrey Series

Infrared

6

7, ...

>75000

When an electric discharge is applied to gaseous hydrogen, the H2​ molecules break apart, forming excited hydrogen atoms. These excited atoms emit electromagnetic radiation at specific, discrete frequencies, leading to the hydrogen emission spectrum, which consists of several spectral lines named after its discoverers.

In 1885, Johann Balmer observed that the visible lines in the hydrogen spectrum could be described using a mathematical formula. When expressed in terms of wavenumber

 (νˉ), these lines follow:

νˉ=RH​(221​ - n21​)

where:

  • RH is the Rydberg constant,
  • n is an integer equal to or greater than 3 (i.e., n = 3,4,5,……).

This equation specifically applies to the Balmer series, which corresponds to the visible region of the hydrogen spectrum.

Table of Contents


  • 1.0Spectral Emission
  • 2.0Spectral Series of Hydrogen

Frequently Asked Questions

The hydrogen emission spectrum is the pattern of light emitted by hydrogen atoms when they are excited and release energy. The spectrum is composed of distinct lines, each corresponding to a specific wavelength of light. These lines are grouped into series, such as Lyman, Balmer, Paschen, Brackett, and Pfund, based on the electron transitions within the atom.

The discrete lines in the hydrogen emission spectrum are due to the quantised energy levels in hydrogen atoms. When an electron transitions from a higher energy level to a lower one, it releases energy as a photon with a specific wavelength, leading to a distinct spectral line.

The discrete lines in the hydrogen emission spectrum are due to the quantised energy levels in hydrogen atoms. When an electron transitions from a higher energy level to a lower one, it releases energy as a photon with a specific wavelength, leading to a distinct spectral line.

Hydrogen has only one electron, making its electron transitions straightforward to track. In contrast, heavier elements have multiple electrons, leading to more complex interactions and, consequently, more complex spectra.

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